Introduction to Acid Base Equilibria
Acid base equilibria describe the dynamic balance between proton donors and acceptors in solution. Central to aqueous chemistry, these equilibria govern pH, reactivity, and biochemical processes. Equilibrium involves reversible proton transfer reactions with defined constants.
"Understanding acid-base equilibria is essential for mastering chemical reactivity and environmental systems." -- Linus Pauling
Definitions: Acids and Bases
Arrhenius Definition
Acid: increases H+ concentration in aqueous solution. Base: increases OH− concentration. Limited to aqueous media.
Brønsted-Lowry Definition
Acid: proton donor. Base: proton acceptor. Applicable in non-aqueous solvents and gas phase.
Lewis Definition
Acid: electron pair acceptor. Base: electron pair donor. Broadest definition, includes coordinate bond formation.
Equilibrium Constants in Acid Base Systems
Acid Dissociation Constant (Ka)
Defines strength: Ka = [H+][A−]/[HA]. Larger Ka indicates stronger acid.
Base Dissociation Constant (Kb)
Defines base strength: Kb = [BH+][OH−]/[B]. Related to Ka by KaKb = Kw.
Water Ionization Constant (Kw)
Autoionization: H2O ⇌ H+ + OH−. Kw = [H+][OH−] = 1.0 × 10−14 at 25°C.
| Constant | Expression | Typical Value (25°C) |
|---|---|---|
| Ka | [H+][A−]/[HA] | Varies (10−1 to 10−14) |
| Kb | [BH+][OH−]/[B] | Varies (10−1 to 10−14) |
| Kw | [H+][OH−] | 1.0 × 10−14 |
pH and pOH: Quantifying Acidity and Basicity
Definition of pH
pH = −log[H+]. Measures acidity; pH < 7 acidic, pH = 7 neutral, pH > 7 basic at 25°C.
Definition of pOH
pOH = −log[OH−]. Complements pH: pH + pOH = 14 at 25°C.
Calculating pH in Various Solutions
Strong acid: pH = −log[acid]. Weak acid: use Ka and ICE table. Similar approach for bases using Kb.
For weak acid HA:HA ⇌ H⁺ + A⁻Initial: c 0 0Change: −x +x +xEquilibrium: c−x x xKₐ = x² / (c−x)Solve for x, then pH = −log xBuffer Solutions and Their Mechanism
Definition and Components
Buffers: solutions resisting pH change upon addition of acid/base. Composed of weak acid and its conjugate base or vice versa.
Henderson-Hasselbalch Equation
pH = pKa + log([A−]/[HA]). Enables pH calculation from component concentrations.
Buffer Capacity
Measure of resistance to pH change. Maximum when [A−] ≈ [HA].
Buffer reaction:HA + OH⁻ → A⁻ + H₂OA⁻ + H⁺ → HAEffect: neutralizes added acid/base, stabilizing pH| Buffer System | Weak Acid | Conjugate Base | Useful pH Range |
|---|---|---|---|
| Acetic Acid/Acetate | CH3COOH | CH3COO− | 4.8 - 6.8 |
| Carbonic Acid/Bicarbonate | H2CO3 | HCO3− | 6.1 - 8.1 |
Conjugate Acid-Base Pairs
Concept and Examples
Acid loses proton → conjugate base. Base gains proton → conjugate acid. Example: HCl/Cl−, NH3/NH4+.
Relation Between Ka and Kb
KaKb = Kw. Strong acid’s conjugate base is weak; strong base’s conjugate acid is weak.
Role in Buffer Systems
Buffers depend on conjugate pairs to reversibly accept/donate protons, stabilizing pH.
Acid-Base Titration Curves
Strong Acid-Strong Base Titration
pH starts low, rises sharply near equivalence point (~7), then plateaus high. Equivalence point: moles acid = moles base.
Weak Acid-Strong Base Titration
Initial pH higher due to incomplete dissociation. Equivalence point pH > 7 due to conjugate base hydrolysis.
Indicators and Endpoint Determination
Indicators change color near equivalence point. Selection depends on expected pH range.
Titration reaction:HA + OH⁻ → A⁻ + H₂OAt equivalence: moles HA = moles OH⁻Calculate pH from hydrolysis of A⁻Le Chatelier's Principle in Acid Base Equilibria
Effect of Concentration Changes
Adding H+ shifts equilibrium left, reducing ionization of weak acid. Removing products shifts equilibrium right.
Effect of Temperature
Endothermic dissociation increases Ka with temperature. Exothermic reactions decrease Ka.
Effect of Pressure and Volume
Minimal effect in aqueous solutions due to constant solvent volume.
Salt Hydrolysis and pH of Salt Solutions
Definition of Hydrolysis
Salt ions react with water, producing acidic or basic solutions depending on ion nature.
Acidic Salt Solutions
Salts of weak base and strong acid produce acidic solutions (e.g., NH4Cl).
Basic Salt Solutions
Salts of weak acid and strong base produce basic solutions (e.g., NaCH3COO).
| Salt Type | Example | Resulting pH |
|---|---|---|
| Strong acid + strong base | NaCl | Neutral (~7) |
| Strong acid + weak base | NH4Cl | Acidic (<7) |
| Weak acid + strong base | NaCH3COO | Basic (>7) |
Polyprotic Acids and Stepwise Equilibria
Definition and Examples
Acids with multiple ionizable protons: H2SO4, H3PO4. Ionize in steps.
Stepwise Dissociation Constants
Ka1 > Ka2 > Ka3 due to increasing difficulty removing protons.
Calculations and pH Determination
Approximate pH from first dissociation for dilute solutions. Consider sequential equilibria for accurate results.
Example:H₂SO₄ ⇌ H⁺ + HSO₄⁻ (Kₐ₁ very large)HSO₄⁻ ⇌ H⁺ + SO₄²⁻ (Kₐ₂ ≈ 1.2 × 10⁻²)Use successive ICE tables for calculationsQuantitative Calculations in Acid Base Equilibria
ICE Tables Method
Initial, Change, Equilibrium concentrations used to solve for unknowns in equilibrium expressions.
Approximations for Weak Acids/Bases
If Ka or Kb << initial concentration, x << c approximation simplifies calculations.
Titration Curve Data Analysis
Use stoichiometry and equilibrium constants to determine pH at various titration points.
Applications of Acid Base Equilibria
Biological Systems
Maintain physiological pH (7.4) via bicarbonate buffer. Enzyme activity dependent on pH.
Industrial Processes
Control of pH in chemical manufacturing, pharmaceuticals, wastewater treatment.
Analytical Chemistry
Titrations for concentration determination. pH indicators and sensors.
Environmental Chemistry
Acid rain impact, ocean acidification, soil pH management.
References
- P. Atkins, J. de Paula, Physical Chemistry, 10th ed., Oxford University Press, 2014, pp. 678-723.
- M. Brown, T. LeMay, B. Bursten, Chemistry: The Central Science, 14th ed., Pearson, 2017, pp. 605-660.
- R. Chang, General Chemistry: The Essential Concepts, 7th ed., McGraw-Hill, 2010, pp. 425-470.
- J. McMurry, R. Fay, General Chemistry, 5th ed., Pearson, 2012, pp. 510-558.
- S. Silberberg, Chemistry: The Molecular Nature of Matter and Change, 7th ed., McGraw-Hill, 2013, pp. 455-500.